1334
Scom, DOV~LIS, I~ISKC, HVEBED, MESSERLI , ROSSENLOPP, ISD XCCVLLOUGH Vol. 66
2-METHYL-%BUTANETHIOL: CHEMICAL THERMODUNANIC PROPERTIES AND ROTATIOKAL ISOMERISM1 BY D. M'. SCOTT,D. R. DOUSLIN,H. L. FINKE, TV. K. HUBBARD, J. F. MESSERLY, I. A. HOSSER'LOPP, AND
J. P. IICCULLOUGH
Contribution N o . l i s from the Thermodynamics Laboratory of the Bartlesville Petroleum Research Center, Bureau of Mines, U.S. Department of the Interior, Bartlesville, Okla. Received February 6 , 1862
Thermodynamic and spectroscopic properties of 2-methyl-2-butanethiol (t-amyl mercaptan) were measured, and t h e results were used for calculating the chemical thermodynamic properties in the ideal gas state (0 to 1000°K.) and for studying the rotational isomerism of the compound. Experimental thermodynamic studies provided the following information: values of heat capacity for the eolid (12°K. to the triple point), the liquid (triple point t o 347"K.), and the vapor (360 to 500°K.). the triple point and transition temperatures; the heats of fusion and transition; thermodynamic functions for the solid and liquid (0 to 350°K.); heat of vaporization (330 to 372°K.); parameters of the equation of state; vapor pressure (324 to 411'K.); and the standard heat of formation a t 298.15"K. The infrared spectrum of the crystals and relative intensity U S . temperature data for a pair of infrared bands of the liquid were determined.
Thermodynamic studies of 2-methyl-2-butanethiol (t-amyl mercaptan)
temperature infrared spectrum of the crystals from this research, are collected in Table I. Two moderately intense Raman and infrared bands of the CH, CH, liquid, 577 and 628 cm.-1, appear in the region in CH3-&--CH2 / which only C-S stretching fundamentals are exI pected. Arguments for assigning the lower freSH quency to the C1 conformations (which are optical were made as part of continuing research on organic isomers and have identical spectra) and the higher sulfur compounds, The experimental work con- frequency to the C, conformation have been given sisted of low temperature calorimetry, vapor flow by Brown and Sheppardj for the analogous t-amyl calorimetry, combustion calorimetry, and compara- chloride. As only the lower frequency persists in tive ebulliometry, supplemented with special infra- the infrared spectrum of the crystals, 2-methyl-2red spectroscopic studies. Experimental values of butanethiol crystallizes as the C1 conformation, the entropy and heat capacity in the ideal gas state, and all frequencies observed in the spectrum of the obtained as described in the Experimental section, crystals must be assigned to the C1 conformation. were interpreted by methods of statistical mechanics. The spectrum of the crystals in the sodium chloride The results were used to calculate a table of thermo- and potassium bromide regions, 400 to 5000 cm. -l, dynamic functions and, with an experimental value differs i'rom that of the liquid by the absence of a of the heat of formation, also the heat, free energy, few bands peculiar to the C, conformation and by a and equilibrium constant of formation for tem- general sharpening of the bands and improved resoperatures of interest. The energy difference be- lution. However, these differences are not very tween the rotational isomers of the compound was striking. I n the cesium bromide region, 285 to determined from the spectroscopic studies. 400 cm.-l, on the other hand, the differences between the spectra of crystals and liquid are so Analysis by Methods of Statistical Mechanics.--The analysis by methods of statistical mechanics striking that the spectra are scarcely recognizable will be discussed first. ,411 54 degrees of freedom as belonging to the same substance. In the specof the molecule had to be accounted for. These trum of the crystals, each band of the liquid is degrees of freedom may be classified as 3 transla- split into two components by interactions with the tions, 3 over-all rotations, 43 vibrations, and 5 crystal lattice, and these components are relatively internal rotations. One of the internal rotations, sharp and intense, whereas the corresponding bands that of the ethyl group relative to the rest of the of the liquid are broad and weak. One fundamolecule, involves three stable conformatioils or mental not even detected in the spectra of the rotational isomers, as illustrated in Fig. 1. liquid appears plainly as the doublet 336-348 cm.-I (Strictly, rotation of the thiol group also involves in the spectrum of the crystals. rotational isomers but, because of the small mass of The description of the vibrational modes in the off-axis hydrogen atom, these isomers do not Table I is somewhat schematic and is intended differ significantly in their spectroscopic or thermo- merely to show that the expected number of fredvnamic properties and need not be considered.) quencies is assigned in each region of the spectrum. The most direct evidence for rotational isomerism The CH, and CH2bending and C-€I stretching frein 2-methyl-2-butanethiol comes from interpreta- quencies are not all resolved, and average values tion ol' the molecular spectra, which will be con- had to be selected for thermodynamic calculatims. sidered next. Two frequencies, a C",rocking and a C -C stretchSpectra and Interpretation.-The molecular (3) American Petroleum Institute Research Project 44 a t the Carspectra from various sources,2-4 including a low negie Institute of Technology, Catalog of Raman Spectral Data, serlal (1) This iniestieation was part of 4mrrican Petroleum Institute Research Project 48.4 on the "Production, Isolation a n d Purification of Sulfur Compounds and Measurement of Their Properties," u hich the Bureau of Mines conducts a t Bartles\1lle, Okla., and Laramle 11 yo. (2) H.W, F. Kohlrauseh and F, Koppl, Mmatah+,B P , 255 (1833)s
N o 350. (4) Ref 3, Catalog of Infrared Spectral Data, Serlal No 1723 1724 and 2194 ( 5 ) J K E r o u n and N Shrgpard, Xruns Faraday S a c , 60, 1164 (1068)d
July, 1962
THERhIODPNAJIIC PROPERTIES O F 2-;\I~THPI.-2-BCTllsETIlIOL
1335
TABLE I hf OLECULAR SPECTRA OF Raman (with intensity*), liquid liquid (ref. 2) (ref. 3)
2-lCIETBPL-2-BUTANETHIOL IN
CM.-'
AND INTERPRETATION
Interpretation (description of modea for GI fundamentals)
Infrared (with intensitya) crystals I1 ( 1909
(3%
-
Not investigated
C-C-C bending
232 (lb.)
236 (2.5)
298 (1)
300 (3.5)
299 m-
C-C-S bending
326 ( I )
(332)b
324 vw
C-C-S bending C-C-C bending
348 m 369 w
C-C-C bending
(393)b
391 TV
C-C-C bending
413 (2) 474 ( 1 ) 575 (10) 027 (3) 776 (3)
414 480 576 629 778
808 (3)
808 (9.7)
414 vw 476 w 5X s 627 m 778 sh 787 w 810 s
370 (4)
370 (11.4)
(4.4) (2.4) (40.7) (11.8) (5.1)
860 (2)
860 (6.0)
864 m
923 (2sb)
920 (5.0)
922 m
1011 (2.7) 1060 ( 4 . 5 ) 1139) ( 4 . 9 ) 1163 (7.6) ( l184)b
942 sh 999 sh 1013 m 1063 m 1142 s I107 s 1190 sh
1220 (3.5)
1222 w
1288 (2.6)
1292 in 1342 sh 1379 s
(941)b 1007 (2sb) 1057 (2) 1135 (2) I160 (2) 1215 (1) 1284 ( I )
( 1334)b ( 1385)b
1442 (6b)
478 m 573 s 772 m 806 s ea. 867sh
\
m 1 { 918 929 m ca. 937 sh
1003 s 1062 m
+ +
1
1170 s ea. 1194sh / 1215 w \ 1225 rn 1285 m 1337 w 1376 m
1441 (16.1) 1464 8
Other isomer 2 X 236 = 472 C-S stretching Other isomer CHr rocking Other isomer C-C stretching C-S-H bending & 300 577 = 877 CHs rocking & 342 577 = 919 CH3 rocking Other isomer CHI rocking CHa rocking CHI rocking & 1 2 X 577 = 1154 328 862 = 1190 C-C stretching G C stretching CH2 wagging CH2 twisting CHs sym. bending CHg unsym. bending & CH? bending
1451 in
+
Region 1500-2500 cm.-l omitted S-H stretching 2575 (34.4) 2591 2520 m 2889 (38.6) 2903 (48.1.) 2925 (54) C-H stretching 2924 s 2880 s } 2937 (52) I 2967 (8sb) 2967 (54) J Infrared intensities are given by the abbreviations: s, strong; a Raman intensities are given as in the original reference. Parentheses denote very weak Raman bands read off m, medium: w, weak; vw, very weak ; sh, shoulder; ca., about. of the published tracing. 2571 (3b) 2884 (3b) 2898 ( 5 ) 2923 (6b)
ing according to the schematic description, cannot be assigned from the obserred slpectra. As will be discussed later, the two unobserved frequencies were given an average value of 1290 cm.-l to fit the calorimetric data. The complete set of frequencies selected for the C1 conformation, with the two unobserved values in brackets, is : skeletal bending, 236, 300, 328, 342, 370, and 392; C-S stretching, 577; CHZrocking, 778; C-S-H bending, 86%; C-C stretching, 809, 1221 ( 2 ) , and [1290]; CH3 rocking, 921, 942, 1012, 1062, 1153, and [12901; CHz wagging, 1290; CHz twisting, 1342; CH, and CH,
1
bending, 1380 (3) and 1450 (7); S-H stretching, 2575; and C-H stretching, 2950 (11) cm.-*. Moments of Inertia and Internal Rotation.Valt~esof moments of inertia needed for treating the contributions of over-all and internal rotation to the thermodynamic functions were computed for the GI conformation with the thiol hydrogen atom in one of the three stable positions. Bond distances and acgles were takeq to be: C-C, 1.:54 PI.; C-H, 1.09 A.; C-S, 1.819 A.; S-H, 1.336 A.; C-S-H angle, 96" 30', all other angles tetrahedral. These values are partly the usual ones for paraffin
1336
SCOTT,DOUSLIN, FINKE, HUBBARD, MESSERLY, HOSSENLOPP, AND NCCULLOUGH Vol. GG I
C, Fig. 1.-Rotational
CS
CI
isomers of 2-methyl-2-butanethiol.
hydrocarbons; the rest are transferred from methanethiol.6 The calculations were made by the general method of Kilpatrick and Pitzer.7 The product of the principal moments of inertia for the assumed structure is 3.966 X 10-11* 8 . 3 cm.6. The “effective” reduced moments of inertia, taken so their product over all internal rotations equals [D], the determinant of the internal rotational kinetic energymatrix, are 38.06 X 10-40and5.068 X g. cm.2 for the ethyl group and its methyl group, respectively; 5.223 X and 5.202 X g. cm.2for the other two methyl groups; and 2.833 X g. cm.2 for the thiol group. The symmetry number is 3 for each of the methyl rotations and one for the other internal rotations and for over-all rotation. Simple threefold cosine-type barriers to internal rotation were assumed for the thiol and methyl groups. Values of the barrier height were taken to be 1.5 kcal. mole-1 for the thiol group, as in other alkanethiolss; 3.4 kcal. mole-1 for the methyl group of the ethyl group, as in propaneg; and 5.0 kcal. mole-‘ for the other two methyl groups, the average of the values 5.1 kcal. mole-’in 2-methyl-2propanethiol1° and 4.9 kcal. mole-l in 3,3-dimethyl2-thiabutane.” The analogy to 2-methyl-2-propanethiol and 3,3-dimethyl-2-thiabutanealso was used to estimate the height of the lower barrier for rotation of the ethyl group as 5.0 kcal. mole-’. For the energy difference between the C, and C1 conformations, the observed value of 0.25 ( 1 0 . 2 5 ) kcal. mole-l, determined from spectroscopic data as described in the Experimental section, was used. This value is reasonable in terms of the steric interactions involved. The C1 conformation involves one gauche n-butane interaction and one gauche 1-propanethiol interaction, whereas the C, conformation involves ttvo gauche n-butane interactions. A rough approximation to the energy difference would be the difference between a gauche n-butane interaction, 0.8 kcal. mo1e-1,*2 (6) T. Kojima and T. Nishikawa, J . Phys. Soc. Japan, 12, 680 (1967). ( 7 ) J. E. Kilpatrick and K. 8. Pitzer, J . Chem. Phya., 17, 1064 (1949). (8) J. P. McCullough, H. L. Finke, D. W. Scott, M. A. Gross, J . F. Messerly, R. E. Pennington, and G. Waddington, J . A m . Chem. Soc., 76, 4796 (1954). (9) K. S. Pitzer, J . Chem. Phyr., l a , 310 (1944). (10) J. P. McCullough, D. W. Scott, H. L. Finke, W. N. Hubbard, M. E. Gross, C. Katz. R. E. Pennington, J. F. Messerly, and G. Waddington, J . Am. Chem. Soo., 7 6 , 1818 (1963). (11) D. W. Scott, W. D. Good S. 9. Todd, J. F. Messerly, W. T. Berg, I. A. Hosaenlopp, J. L. Lacina, A. Osborn, and J. P. MeCullough, J . Chem. Phys., S6, 406 (1962). (12) S Mizushima, “Structure of Molecules and Internal Rotation,” Academic Press, Ino., R’aw York, N. Y., 1964, p. 98.
0
2ll
@ Fig. 2.-Potential energy function for internal rotation of the ethyl group in 2-methyl-2-butanethiol.
and a gauche 1-propanethiol interaction, 0.4 kcal. m01e-l.’~ This rough approximation to the energy difference, 0.4 kcal. mole-1, lies within the limits set by the experimental observations. The complete potential energy function assumed for the ethyl rotation is illustrated in Fig. 2. The contributions to the thermodynamic functions were obtained from recently published tables for potential energy functions of this type.14 With the assumed potential barriers to internal rotation, all contributions to the entropy and heat capacity could be calculated except those of ttvo vibrations and of vibrational anharmonicity. When an average value of 1290 cm.-1 was taken for the two unobserved frequencies and no correction was made for vibrational anharmonicity, excellent agreement was obtained with the experimental values of entropy and heat capacity, as shown in Table 11. As 1290 cm.-l is a reasonable value for the average of the two unobserved frequencies, the excellent agreement with the calorimetric data is evidence of the general correctness of the assumed potential barriers to internal rotation. “FABLE 11 OBSERVED AND CALCULATED THERMODYNAMIC PROPERTIES O F 2-METHYL-2-BUTANETHIOL ?’,OK.
330.22 349.77 372 29
Entropy, So, eel. deg.-l mole-1 Obsd. Caled.
96.11 98.28 100.73
96.12 98.30 100.77
Heat cal. deg.-l eapaoity, mole-’ Cp0, T,OK.
Obsd.
Calcd.
380.25 381.20 410.20 451.20 500.20
39.54 41.25 43.59 46.75 50.31
39.54 41.27 43.80 46.76 50.29
Chemical Thermodynamic Properties.-The molecular parameters discussed in the previous sections were used in calculating the thermodynamic functions of 2-methyl-2-butanethiol for (13) R. E. Pennington, D. W. Scott, H. L. Finke, J. P. McCullough, J. F. Messerly, I. A. Hossenlopp, and G. Waddington, J . Am. Chem. Soc., 78, 3266 (1956).
(14) D. W. Scott and J. P. McCullough, Bureau of Mines Report of Investigation KO.5930 (1901).
THERMODYNAMIC PROPERTIES
July, 1962
TABLE I11 THEMOLALTHERMODYNAMIC PROPERTIES OF %METHYL-%BTJ%'ANEl'EIOL 1'. OK.
(8'"
- Hoo)/T, (He-H 0 o ) / T , H o
Gal. dee. -1
oal. deg, "1
- H'o,
kcal.
BO.
cal. dag. -1
1337
OF 2-h!fETHYL-2-BUTA4NETHIOL
I N TEE IDEAL GAS AFf" b
AHfo,b
CPO,
cal. deg. -1
kaal.
lscai.
STATE' log K f
b
-37.59 Infinite 0 0 -37.59 0 0 0 -10.53 8.42 -45.21 32.21 89.56 20.07 6.483 69.49 - 7.33 5.38 -45.75 34.30 6,316 92.48 21 18 71.30 5.17 7.10 -45.79 34.46 92,69 21.26 6.379 71.43 -j- 6.10 - 3.33 -47.73 103.75 42.79 10.24 78.15 25.60 8.63 -49.26 19.74 50.28 114.13 14.91 29.81 84.32 33.66 -12.26 -50.42 56.58 123.86 20.26 90.10 33.76 600 47.74 -14.91 -51.25 61.84 133.00 26.19 700 95.59 37.41 61.92 -16.91 -51.81 66.28 141.56 32.60 40.75 800 100.81 76.16 -18.49 -52.13 70.05 149.58 39.42 105.79 43.79 900 -52.23 90.43 -19.76 73.30 46.59 157.14 1000 110.55 46 59 a To retain internal consistency, some values in tJhis table are given to one more decimal place than is justified by the The standard heat, standard free energy, and common logarithm of the equilibrium constant for the absolute accuracy. formaftionof 2-methyl-2-butanethiol by the reaction, 5C(c, graphite) 6H4g) 1/2 Sz(g) = C,H,,S(g). selected temperatures up to 1000°K.16 The re- and mass measurements were referred to standard devices calibrated a t the National Bureau of Standards. The sults are in columns 2-6 of Table 111. energy equivalent of the combustion calorimetric system, The calculated values of the thermodynamic &(Calor.), was determined by combustion of benzoic acid functions, the experimental value of AHf O298.16 (NBS Sample 39g certified to evolve 26.4338 f 0.0026 kj. (-45.75 =k 0.23 kcal. mole-I), and values of the per gram mass under specified conditions). Material.-The sample of 2-methyl-2-butanethiol was thermodynamic functions of C(c, graphite),16 part of the Standard Sample of Organic Sulfur Compound Hz(g),16and S2(g)17were used in computing values API-USBM 26, prepared at the Laramie (Wyo.) Petroleum of AH?, AFf', and log Kf. The results are in Research Center of the Bureau of Mines.*' The purity, determined by calorimetric studies of melting point aa a columns 7-9 of Table 111. function of fraction melted, was 99.89 f 0.05 mole yo. Heat Capacity in the Solid and Liquid States.-The obExperimental The basic experimental techniques used for 2-methyl-2- served values of heat capacity, C,, are listed in Table IV. butanethiol are described in published accounts of apparatus Above 30"K., the accuracy uncertainty is estimated to be and methods for low temperature calorimetry,18 vapor flow no greater than 0.2%. The crystals undergo an isothermal calorimetry,] comparative ebulliometry,20 and combustion transition about 10" below the melting point, and the low calorimetry.*1 The reported values are based on a molecular temperature crystals (Crystals 11) have a Mike anomaly weight of 104.212 g. mole-' (1951 International Atomic in the heat capacity that reaches a maximum a t 145°K. WeightP), the 1951 values of the fundamental physical Hysteresis in the thermal behavior was observed in the temconstants,zSand the relations: 0" = 273.15OK.24 and 1cal. = perature region of the anomaly. The heat capacity of the 4.184 j .. (exactly). Measurements of temperature were liquid varies regularly with temperature and may be repremade with platinum resistance thermometers calibrated sented within 0.03% between 200 and 350°K. by the empiriin terms of the International Temperature Scale25 between cal equation 90 and 500°K. and the provisional scalez6of the National C,(liq) = 53.373 0.140932' 5.425 X 10-4TT" 4.6094 X Bureau of Standards between 11 an 90°K. All electrical 10-7T8 cal. deg.-l mole-' (1) 0 273.15 298.15 300 400 500
-
I
-
I
*
+
+
-
+
-
_ _ I -
(15) The vibrational contributions were computed b y the Bureau of Mines Electronic Computer Service, Pittsburgh, Pa., and no corrections for vibrational anharmonicity were applied; the contributions of the methyl and thiol internal rotations were computed by Denver Electronic Computing Service, Inc., by two-way curvilinear interpolation in the tables of X. S. Pitzer and W. D. Gwinn, J . Cham. Phys., 10, 428 (1942). (16) D. D. Wagman, J. E. Kilpatrick, W. J. Taylor, K. 8. Pitzer, and F. D. Rosoini, J . Res. Natl. Bur. Std., 34, 143 (1945). (17) W. H. Evans and D. D. Wagman, ibid., 49, 141 (1952). (18) H. M. Huffman, Chem. Rev., 40, 1 (1947); H. M. Huffman, S. S. Todd, and 6. D. Oliver, J . Am. Cham. SOC., 71,584 (1949): D. W. Scott, D. R Douslin, M. E. Gross, G. D. Oliver, and H. M. Huffman, %bid.,74,883 (1952). (19) G. Waddington, S. S. Todd, and H. M. Huffman, ibid., 69, 22 (1947); J. P. 1McCullough. D. W. Scott, R. E. Pennington, 1. A. Hossenlopp, and G. Waddington, ibid., 76, 4791 (1954). (20) 0. Waddington, J. W. Knowlton, D. W. Scott, G. D. Oliver, S. El. Todd, W. N. Hubbard, J. C. Smith, and H. M. Huffman, ihid., 71, 797 (1949). 121) W. N. Hubbard, C. Katz, and G. Waddington, J . Phys. Chem., 68, 142 (1954). (22) E. Wichers, J . A m . Chem. Sac., 74, 2447 (1952). (23) F. D. Rossini, F. T. Gucker, Jr., H. L. Johnston, LoPauling, aud G. W. Vinal, ibid., 74,2699 (1952). (24) Some of the results originally were computed with constants and temperatures in terms of the relation 0' =: 273.16OK. Only results affected signifioantly by the new definition of the absolute temperature scale [H. F. Stimson, A m . J . Phye., a3, 814 (1955)l were recalculated. Therefore, numerical inconsistencies, much smaller than the accuracy uncertainty, msy be noted in some of the reported data. (25) H. F.Stimson, J . Res. Natl. Bur. Std., 42, 209 (1949). (26) H.J. Noge and F. G. Brickwedde, ibid., 22, 851 (193Q).
Heats and Temperatures of Fusion and Transition.-Two determinations of the heat of fusion, AHm, gave the values 145.25 and 145.51 cal. mole-'; the accepted value is 145.4 i 0.2 cal. mole-'. Three determinations of the heat of transition, AHt, gave the values 1907.3, 1907.0, and 1907.0 cal. mole-I; the accepted value is 1907.1 f 0.2 cal. mole-'. 2-Methyl-2-butanethiol i s seen to belong to the class of substances with globular molecules for which the heat of transition is large and the heat of fusion small. The energy associated with the anomaly a t 145°K. is only about 8 cal, mole-'. The results of two separate studies of the melting temperature, TF,aa a function of fraction of total sample melted, F, are reported in Table V. These results, which indicate that the impurity forms a solid solution with the major component, were treated by the method of Mastran elo and Dornte.** Listed in Table V are the values obtainef for the triple point tem erature, Tt,; the mole fraction of impurity in the sample, 82*, the Henry's law constant for distribution of the impurity between the solid and liquid phases, K ; and the cryoscopic c0nstants,~9A = AHm/RTt,z and B = 1/Tt, - ACm/ZAHm, calculated from the observed values of Tt,, AHm, and ACm (1.28 cal. deg.-lmole-l). The temperature for equilibrium between crystals I and I1 was studied as a function of per cent of total sample in the high temperature form (Crystals I) with the results: 20.42y0, (27) J. C. Morris, W. J. Lanum, R. V. Helm, W. E. Haines, G. L. Cook, a n d John 8. Ball, J . Chem. Ens. Data, 6 , 112 (1960). (28) 8. V. R . Mastrangelo and R. W. Dornte, J . A m . Chem. BOG.,71, 6200 (1955). (29) A. R. Glasgow, A. J. Streiff, a n d F. D. Rossini, J . Res. Natl. Bur. Std., 36, 355 (1945).
1338
SCOTT, DOUSLIX, VIXKE, I I U B B A R D , I\IESSERLP,
HOSSEKLOPP, bND
&ICCULLOCGI-I
Vol. 66
TARLE IT'
MOLALHEAT CAPACITY X, OK,a Cs 0 Crystals 11
T.
OF 2-?IIETHYL-z-BTTA'\rETHIOL IN G A L .
ATb
CB
T,OK.d
DE,E.-~ CB C
ATb
100.57 5.551 18.791 149.47 32.4 85# 3.940 102.02 5,379 19.060 149.92 33.02P 2.057 12.43 103.01 I . 075 5.401 19,226 151 92 35,161" 1,956 12.95 104.96 1.199 6.102 19.566 2.222 152.54 36.13fP 13.69 1.398 106.00 5 317 38. 624d'E 153.80 19.747 1.815 14.50 1.620 110.93 5.821) 20.633 15.06 116.64 3.686 1.786 21. 715 Crystals I 16.02 117.23 2.062 5,896 21.835 122.11 16 47 2.204 5.362 22.816 40.696r"*e 160.73 2.083 17.57 2.532 5.318 122.66 22.943 4 0 . 78gd*# 162.34 1.075 17.96 2,653 127.37 5.151 23.972' 162.81 40. 93id8" 2.073 19.30 3.062 127.87 5.106 24.110e 40. 967d,8 163.88 2.003 19,67 4.879 3.170 131.19 24,917" 41, 02Sd?" 164.87 2.066 21.17 3.635 131.41 4.966 24. 96gg 41. 065d*8 1,995 165.87 21.58 5.909 132.90 3.763 2 5 . 34lS 23.23 4.236 135.97 4.675 26,279' Liquid 23.73 4.397 4.755 136.27 26,346' 26.23 5.119 137.37 3.388 42.671 26.711" 171.53 3.905 26.32 5.582 5.147 138.64 42,753 27.213' 176.72 6.481 28.69 2.042 140.09 5.778 27.811" 42,809 10.501 179.50 31.47 4.533 140.79 6.498 28,144" 42,858 6.453 183.19 34.74 2.045 7.249 140.96 28.209* 42,982 180.62 6.419 2.047 141.69 38.56 8.034 42,984 28,185" 10.420 189.96 42.83 142.09 1.983 8.861 43.206 28,900" 200.54 10.404 142.89 47.48 1,952 9.705 43.454 29.51ie 210.89 10.317 53.24 143.27 1.126 43.744 10.703 28.926" 10.229 221.16 54.10 144.02 1.896 44.096 231.34 10.841 10.134 30. 543e 144.31 54.27 44.466 10.881 953 10.037 241.43 29. 384e 144. 80 44 922 57.34 1.875 11.398 10.922 251.91 31. 0338 144.88 59.81 1.901 45.436 11.816 262.77 10.804 31.012" 144.95 60.00 45.990 11.849 1.399 10.667 29.641b 273.51 145.25 61.74 46.039 12.155 9.608 29,978" ,938 274 I54 65.74 145.41 46.332 10.365 12.828 31,176' 280.21 ,901 145.79 46.921 71.48 1.940 10.238 13.768 30.208" 290.51 146.17 10.116 47.508 77.25 ,924 14.740 30.550" 300.69 47.552 1.371 146.32 83.04 9.326 15.769 301.23 30. 4066 47,800 2.148 146.79 88.91 9.996 16.790 305.73 30.913" 48.164 146.92 10,126 2.145 89.19 16.858 310.96 30,951" 147.70 48.452 2.147 9.872 91.67 315.67 31,285" 17.258 49,118 10,636 325.92 147.80 94.88 31.348" 1.605 17.807 49.812 10.486 2.145 147.83 96.51 336.48 31. 347e 18.085 50,534 148.82 10.335 346,89 31,930" 2.100 97.48 18.249 b AT is the temperature increment of each heat e T is the mean temperature of each heat capacity measurement. Results for the C, is the heat capacity of the condensed phase at saturation pressure. ca acity measurement. soid are not corrected for the effect of premelting due to the presence of impurities. e Correction applied to mean heat capacity for curvature. I
I
158.79'K.; 46.527& 158.98'K.; and 77.787,, 159.06'K. The transition temperature was taken to be 159.1 f 0.1"K. Thermodynamic Properties in the Solid and Liquid States. -Values of thermodynamic functions for the condensed phases were computed from the calorimetric data for selected temperatures between 10 and 350°K. The results are in Table VI. The values at 10°K. were computed from a Debye function for 5.5 degrees of freedom with 0 = 112.7'; these parameters were evaluated from the heat capacity data between 12 and 20°K. Corrections for the effects of premelting have been applied to the "smoothed" data in Table VI. Vapor Pressure.-Values of vapor pressure determined by comparative ebulliometrg with water as the reference substance are in Table VII. These values were obtained in the second of two attempts t o measure the vapor pressure. The first attempt was abandoned when the compound started to decompose with formation of highly volatile impurity, but the second attempt with a fresh sample was carried through without any detectable decomposition. The different stability of the 2 samples is not understood.
Intermittent tests for decomposition of the second sample were made by repeating measurements at the lowest pressure after each measurement a t a higher pressure. In 14 such replicate determinations, the observed boiling points deviated from the mean at most 0.005' and seldom more than 0.002'. The difference between the ebullition and condensation tpmperatures of the sample at 1 atm. pressure was less than 0.001 ". The Antoine and Cox equations selected to represent the results are log p = 6.82837 - 1284.885/(t 218.750) (2) log ( p / 7 6 0 ) = A ( 1 - 372.280/2') (3) log A = 0.841629 - 7.4895 X 10-4T 7.3565 X 10-'T2 In these equations, p is in mm., t is in "C., and T is in "K. Observed and calculated vapor preesurev for both equations are compared in Table VII. The normal boiling point calrul a t d from either equation is 99.13' (372.28"K.). Heat of Vaporization, Vapor Heat Capacity, and Effects of Gas Imperfection.-The experimental values of heat of vaporization and vapor heat capacity are in Tables VI11 and
+
+
July, 1962
IX. The estimated accuracy uncertainty of the values of AHV and Cueare 0.1 and 0.2 respectively. These values
were obtained in the second oT Otwo 3 attempts to do vapor flow calorimetry of 2-methyl-2.butanethiol. The first attempt, like the first attemph to study the vapor pressure, was abendoned because of decomposition of the sample. The sample used for the first attempt was somewhat less pure than the one used for the second, successful attempt, but the decomposition of the first sample was not necessarily related to the impurity. Tests for effects of any decomposition of the second sample were made from time to time during the vapor flow ca.lorimetry. Although evolution of minute traces of HzS mas detected as measurements were made at progressively higher temperatures, no other evidence of decomposition was observed until t,he heat capacity was determined a t the highest temperature, E100.2"K. During the heat capacity measurements at 500.2"K., the vapor pressure of the sample in the vaporizer increased significantly (7 mm. a t 760 mm. total pressure) as if a decomposition product of higher volatility than 2-methyl-2-butanethiol was being
TABLE V
MELTINGP O I X T
%hf:ETHYL-2-BUTANETHIOL:
Tt,
SUMMARY
169.37 rt 0.10'K.; Ns* = 0,0011 f 0.0005 mole fraction; K = 0.221; A = 0.002552 deg.-l; B = 0.001490 deg. -* =
Series
I I I1 I I1 I1 I I1 I I1
(1[' -I-K)1 K'-l
F
0.0999 .1735 .3227 .3843 ,4298 .5834 .6333 .7459 .8650 ,9472 1.0000 Pure
TF,OK.
2.607 2.187 1.649 1.497 1.402 1.153 1,091 ,971 .871 .812 .779
T d o d , OK.
168.260 168.439 168.673 168,729 168.773 168.877 108.901 108.955 168.998 169.032
168.264 168.441 168.668 168. 733 168,773 168.878 168.904 168,955 168.997 169.022 169.036 169.365
TABLEVI THEMOLALTHERMODYNAMIC PROPERTIES O F 2-METHYL-2BUTANETHIOL I N THE
-
(Hs
-
-(Fa Hoo)/T, cal. deg. -1
";a?l!'"'
10 15 20 25 30 35 40 45 50 60 70 80 90 100 110 120 130 140 145 145 150 155 159.1
0.050 ,164 ,366 .649 .996 1.389 1 814 2,260 2.720 3.661 4.617 5.575 6.531 7.485 8.435 9,382 10.327 11.273 11.748 11.748 12.226 12 710 13.112
0.148 ,471 ,982 1.591 2.235 2.877 3,495 4.084 4.646 5.705 6.703 7.660 8.598 9.521 10.435 11.349 12,280 13.262 13.837 13,837 14.417 15.108 15.791
159.1 160 169.3
13.11 13.27 14.87
27.775 27.850 28.671
T, OK.
1339
'rHERMODYNAMIC PROPERTIES O F 2-&fETHYL-2-BUTSSETHIOL
SOLID AND LIQUIDSTATESb Ha
deg. -1
- H'o,
enl.
SS.
eal.
deg. -1
CS,
cal. deg. -1
Crystals I 1.480 7.066 19.630 39.77 67.04 100.70 139.80 183.79 232.31 342.3 469.2 612.8 773.8 952.1 1147.8 1361.9 1596,4 1856.7 2006.4 2006.4 2162.6 2341.8 2512.3
0.198 ,635 1.348 2.240 3.231 4.266 5.309 6.344 7.366 9.366 11.320 13.235 15,129 17.006 18,870 20.731 22.607 24.535 25.585 25,585 26.643 27.818 28.903
0.588 1.764 3.273 4.764 6.123 7.304 8.321 9.262 10,142 11.848 13.524 15.220 16.973 18.690 20.470 22.378 24.614 27.760 37.5 29.74 33.04 38.92 44.20
40.89 41.12 43.44
40.56 40,64 41.36
Crystals I1 4419 4456 4837
Liquid 160,3
L4,87
170
14,09 16.69 18.34 19.94 21.49 23.00 24.46 25.87 27.25 28.5'3 29.00 30.31 31.18 32.41 33.41 33.63 34.82 35.98 37.11 38.23 39.32
180 190
200 210 220 230 240 260 260 270 273.15 280 290 298.15 300 310
320 330 340 350
29.427 29.482 30.22 30.89 31.50 32.06 32.58 33.07 33.54 33.98 34.41 34.82 34.95 35.22 35.62 35.93 36.00 36.38 36.76 37.13 37.50 37.87
4982
44.20
42.64
6012 5440 5869 6299 6733 7168 7607 8049 8496 8946 9402 9547 9862 10,329 10,713
44.47 46.91 49.23 51.44 53.55 55.58 57.53 59.41 61.24 63.00 64.72 66.26 66.40 68.03 69.34 69.63 71.20 72.74 74.24 75.73 77.19
42.65 42.81 42.99 43.20 43.43 43.72 44.05 44.42 44.84 46.31 46.80 46.96 46.33 46.88 47.36 47.48 48.09 48.74 49.38 50.05 50.75
10,800
11,278 11,762 12,253 12.750 13,254
0 The values tabulated are the free energy function, heat content function, heat content, entropy, and heat capacity of the condensed phases a t saturation pressure.
TABLE VI1 J'APOR PRESSCRE O F 2-hfETHYL-2-BIJTANETHIOL
BoXng point, 'C. 2-Methyl-2butanethiol
Water
p(obsd.),a mm.
p(obsd.1
- p(oalod.), mm.
Antoine eq. 2
Coxeq.3
60.000 50.888 149.41 -0.01 0.00 65 56.725 .00 .00 187.57 70 62.625 233.72 .01 .01 75 68.578 289.13 .01 - .01 355.22 .05 .02 80 74,579 85 80.638 433.56 .04 - .01 90 86.749 525.86 .04 - .02 95 92.914 .OO .05 633.99 100 99.132 760.00 - .03 - .04 105 105.401 906.06 - .05 .01 110 111.728 1074.6 - .1 .o 115 118.106 1268.0 - .2 .o 120 124.537 1489.1 - .2 .o 125 131.021 1740.8 .1 .1 130 13'7.559 .4 .o 2026.0 a From the vapor pressure data for water given by N. S. Osborne, H. F. Stimson, and D. C. Ginnings, J. Res. Natl. Bur. Std., 23, 261 (1939).
+
-+ +
-
+
+
+
+
formed. Tests for vapor phase decomposition in the flow calorimeter, made as described previ~usly,~g showed that decomposition was affecting slightly the results at 500.2"K. Afterwards, similar tests showed no detectable decomposition occurred at 381.2"K. A repeat measurement of the heat capacity at 381.2'K. gave results that agreed within 0.1% with those obtained before the high temperature measurements had been made. The results at 500.2"K. in Table I X were obtained after all the other reported results and so are the only2nes possibly affected by decomposition. The value of Cp a t 500.2'K. was not used in selecting the molecular parameters used in calculating thermodynamic functions. Nevertheless, the experimenhl value a t that temperature agrees well enough with the calculated value to indicate that the effect of decomposition was small even under the most severe conditions. The equation selected to represent the heat of vaporization as a function of temperature is AHv = 11052 4.368T - 1.392 X 10-2TZcal. mole-' (330-3'73'K.) (4) The effects of gas imperfection were correlated by the procedure described in an earlier paper. 3') The empirical
-
(30) J. P. McCullough, H. L. Finke, J. F. hlesserly, R. E. Pennington, I. A. Hossenlopp*and G. Waddington, 1.Am. Chem. Soc.. 77, 6119 (1Q6Q.
SCOTT,DOUSLIN, FINKE,HUBBARD,
1340
&fESSERLY, HOSSENrIOrr, AND
TABLE VI11 THEMOLALHEAT OF VAPORIZATION AND SECOND VIRIAL CO~PPICIENT OF ~-METHYI,-~-BUTANETHIOI, 2',
P , atm.
OK.
E (ohed.),
E(cnlcd.),
00.
00."
AHu, cal.
330.21 0.250 8092zk I* -1705 -1722 849.76 ,500 7821 -I 2b -1478 -1479 372.28 1.000 7497 & 25 -1263 -1259 5 Calculated from eq. 5. b Maximum deviation from the mean of three determinations.
TABLE IX THEMOLALVAPORHEATCAPACITY OF 2-METHYL-2-BUTANETHIOL IN CAL. DEG.-~ T,
OK.
Cp(l.000 atm.) Cp(0.500 atm.) C,(O 250 atm.)
CPO
- TB" (obsd.)"
- TB" (calod.)b a
b
-TB"
360.20
381.20
410.20
451.20
600.20
44.171
47.143
50.563
40.028 39.779 39.54 0.92 0.94
42.069 41.628 41.445 41.25 0.76 0.73
43.732 43.59 0.54 0.54
46.848 48.75 0.38 0.36
50.373 50.51 0.24 0.24
= -T(dZB/dT*),
Calculated from eq. 5.
cal. deg.-l mole-' atm.-l.
equation for B, the second virial coefficient in the equation of state, PV = RT(1 B / V ) , is B = 453 261.2 exp(700/T) cc. male-~(330-50O0K.) ( 5 ) "Observed" values of B and - T(dZB/dTz) = lim(aC,/dP)T
-
+
P-0
and those calculated from eq. 5 are compared in Tables VI11 and TX. The heat of vaporization at 298.15% was calculated by extrapolation with eq. 4 (8.51 kcal. mole-'), by using the Clapeyron equation with eq. 3 and 5 (8.52 kcal. mole-') and by use of a thermodynamic network with the thermodynamic functions of Table 111 (8.53 kcal. mole-'). From the last value, selected rn the most reliable, and eq. 5, the standard heat of vaporization was calculated, AHv'~98.16 = 8.54 kcal. mole-'. Entropy in the Ideal Gas State.-The entropy in the ideal gas state a t 1 atm. pressure was calculated as shown in rable X.
TABLE X THEMOLALENTROPY O F 2-METHYL-2-BUTANETHIOL IN THE IDEAL GASSTATEIN CAT,.MOLE-' T,OK.
e
Nliq.) AHU/T S(idea1) R In Pa
- X(rea1)"
330.22
349.77
74.28" 24.50 0.09 -2.76
77.16" 22.36 0.14 -1.38
_
I
-
-
372.29
80. 37b 20.14 0.22 0.00
-
So(obsd.) f 0.20" 96.11 98.28 100.73 a By interpolation in Table VI. b Extrapolated by use of eq. 1. The entropy in the ideal gas stRte less than in the real gas state, calculated from eq. 5. d Entropy of compression, calculated from eq. 3. e Estimated accuracy uncertainty. Heat of Combustion and Formation.-Results of a typical determination of the heat of combustion of 2-methyl-2butanethiol are given in detail in Table XI. The s mbols syd abbreviations are those of Hubbard, Scott, a n J Wsddington,al except as noted. The results of all experiments are summarized in Table XI1 The derived results in Table XI11 were computed by use of the values of AHc", AHv', S., and So for 2-methyl-2butanothiol and literature values of the standard heat of formation of COz(g),sa HzO(liq),l* H~S0&30H~O(liq)
.
MCGULLOUGH
TdBLE
SUXMARY OB
A
Vo!. 66
XI
TYPICAL COMBUSTION CALORIMETRIC
EXPERIMENT WITH
2-hfETHYL-2-BUTANI3TaIOLa
m' (2-methyl-2-butanethiol), g. Ato = ti ti Atorz., deg.
0,77067 1.99550
- -
-7800.63
?$Calor.) ( - A t , ) , cal. &(Cont.)( - A ~ J ? * cnl. AEmn., tal. AE'deoomp. (ENOs " 0 ~ )ral,~ A E , corr. to st. states,' cal. -me AEc'/M (auxiliary oil), cal. -m"' AEc0/M (fuse), cal.
-
+
27.66 1.35 10.94 2.52 533.48 15.26
m' AEco/M (2-methyl-2-butanethiol), cal. g.-' -7264.74 -9426.52 AEc"/M (2-methy1-2-butanethiol), oal. g.-l Auxiliary data: 8(Calor.) = ,3909.11 oal. deg.-l; V 10983.8 c?l. (Bomb) = 0.347 I . ; AEc"/M (auxiliary oil) g.-'; AEc"/M (fuse) = 3923 cal. g.-'; physical properties a t 25' of 2-methyl-2-butanethio1, p = 0.82119 g. rnl.-1,27 ( h E / d P ) p = -0.011 cal. g.-1 atm.-', and Cp = 0.454 cal. Sf(Cont.) (25' tt deg.-l g.-'. b &'(Cont.).(ti 25') Atoor,.). Items 81-85, incl., 87-91, mcl., 93, and 94 of the computation form of ref. 31.
-
+
- +
TABLE XI1 SUMMARY OF RESULTS OF COMBUSTION CALORIMETRY AT 298.15"K." AEc'/M, cd. g.-' -9429.13, -9425.56, -9428.01, -9426.52, -9423.47, -9425.90, -9428.17 Mean and std. dev. -9426.68 =k 0.73 cal. g,-1 -982.37 i 0. lgb kcal. mole-' AEC0298.16 -985.04 f 0.19* kcal. mole-' AHc'z~a.i: a These results apply to the idealized combustion reaction: 75HzO(liq) = 5Codg) CSHI~S(liq) .l9/20zk) HiSol 80H20(1iq). * The uncertalnty interval equal to twice the final "over-all" standard deviation [F. D. Rossini, "Experimental Thermochemistry, " Interscience Publishers, Inc., New York, X. Y., 1956, Chapter 14, pp. 297-3201.
+
+
+
[-212.07 kcal. m0le-~],3~ and Sz(g)'7 and of the standard entropy of graphite,lG hydrogen gas,'&rhombic sulfur,l' and diatomic sulfur gas," all at 298.15'K.
TABLE XI11 MOLALTHERMODYNAMIC PROPERTIES O F FORMATION 2-METHYL-2-BUTANETHlOL AT 298.15'IC." Reference
state Liquid Gas Gas a
OF
Ref. state ASjo, AFf", log of sulfur cal. deg.-f koal. Kf 0.59 -0.43 S(c,rhombic) -38.87 f.0.21 -132.35 2.23 -1.84 S(a,rhombic) -30.33 jz 0.22 -109.21 5.38 -46.75zt0.23 -128.85 -7.33 S2(g)
':{:;
For the reaction:
5C(c, graphite)
4- 6Hz(g)
+
S(c, rhombic) or 1/2 Sz(g)
= C&S
(liq or 9).
Special Infrared Spectroscopic Studies.-The spectrum of crystalline 2-methyl-2-butanethiol a t liquid nitrogen temperature was determined in a low temperature infrared cell similar to the one described by Biirer.84 However, cesium bromide windows were used so that spectra could be observed to 35 p. The spectrum was determined from 2 t o 15 p with a Perkin-Elmer Model 21 spectrometer with sodium chloride optics and from 15 to 35 p with a PerkinElmer Model 112 spectrometer with cesium bromide optics.S6 The energy difference between the rotational isomers was (33) W.D.Good, J. L. Lacina, and J. P. MoCullough, J . Am. Cham.
(31) W. N. Hubbard, D. W. Scott, and G. Waddington, "Experimental Thermochemistry," F. D. Rossini, Editor, Interscience Publishers, h a . , New York, N. Y.,1956, Chapter 6,pp. 75-128. (32) E. J. Prosen, R. 9. Jessup, and B. D. Rossini, J . Res. Natl. BUT. Std., 88, 447 (1944).
Soc., 82,5589 (1960).
(34) T. Btirer, thesis, "Infrarot-Spektren von Cyclanonen," Eidg. Tech. Hochschule, Zilrieh, Switzerland, 1958. (35) C. A Frenzel, D. W. Scott, and J. P. McCullongh, Bureau of Mines Report of Investigation No. 6658 (1960).
July, 1962
EQUIVALENT CONDUCTANCE IN AQUEOUS THALLOUS HYDROXIDE
studied by determining the ratio of intensity of the infrared bands of the liquid a t 577 cm.-l (GI) and 628 om.-’ (c,) as a function of temperature. The low temperature cell alga WBW used in these atudiee of the temperature dependence of the liquid state spectrum. Determinations were made at the temperaturee obtained with solid carbon dioxide and with nearly boiling water as well as a t room temperature. Values found for the ratio of the integrated intensities of the 577 and 628 cm.-’bands are 3.6 at about 195OK., 2.6 a t about 29S°K., and 2.9 at about 368’K. The variation with temperature is scarcely more than the experimental uncertainty
1341
and indioatea that the rotational isomers have nearly the
aame energy. The value taken for the energy difference waa 0.25 + 0.25 kcal. mole-’ with the C1 isomer more stable. This value was determined for the liquid state but, as the rotational isomers must have small and nearly equal dipole moments, the energy difference is nearly the same in the gaseous state.
Acknowledgments.-The assistance of F. R, From and C. A. Frenzel in some of these experiments is gratefully acknowledged.
EQUIVALEKT CONDUCTAKCE AND IONIC ASSOCIATION I N AQUEOUS TE-IALEOlUS HYDROXIDE SOLUTIONS AT 26’ BY W. T. LINDSAY, JR. Chemistry Department, Westinghouse Research Laboratories, Pittsburgh 36,Pennsylvania Receiaed February 8 , 1968
Previous measurements on thallous hydroxide solutions by Ostwald ( 1887) and Hlasko and Salitowna ( 1935) appear to be in error, probably because of carbon dioxide difficulties. By exclusion of COZ, it has been possible to determine the to equivalent conductance a t 25” with reasonable precision from M. The phoreogram is catabatic in this range. Application of the Fuoss-Onsager equation, with an assumed value of 3.0 A. for the ion size arameter, gives 273.0 f 0.3 for A. and 3.0 f 0.3 for K , (95% confidence limits). The A0 value is in good agreement wit[ that obtained from limiting ionic conductance data. The association constant is close to that obtained by Bell and Prue from reaction kinetic measurements, but is about one-half that derived by Bell and George from solubility equilibria and by Bell and Panckhurst from other kinetic measurements.
Introduction Ionic association in thallous hydroxide so1ution.s is of some interest because of the unreasonably small values obtained for the “distance of closest app:roach” when experimental values of the association constant a,reused in the standard Hjerrum treatment.l The discrepancy originally was ascribed to covalent bonding, but n.m.r.2 and R a i ~ n a nspectra ~ . ~ studies seem to indicate that only electrosta. Lic interactions occur. There also are discrepancies among values of the association constan t obtained by various experimental methods. The reaction-kinetic measurements of B’ell and Prue5 gave 2.63, solubility equilibria results of Bell and George6 gave 6.7, and other reactionkinetic experiments by Bell and Panckhurst’ gavo 7.1, all at 25”. Conductivity data can be used to determine the association constant,, but previous measurements on thallous hydroxide by Ostwalds and Hlasko and Salitownag appear to be in error. As shown in Fig. 1, the trend of t,he Ostwald data is symptomatic of carbon dioxide interference. Although (1) C. W. Davies, Discum-ion? Faraday Soc., 24, 83 (1957): C. W. Davies in “The Structure of Electrolytic Solutions,” W. J. Hamer, Ed., John Wiley and Sons. New York, N. Y.,1959, p. 81; R. P.Bell and hl. H. Panckhurst, Rec. trav. chim., 76, 725 (1956). (2) R. Freeman, R . P. H. Gasser, R. E. Richards. and D. H. ’Xheeler, Mol. Phys.. 2 , 7 5 (lQ59); R. E’. H.Gasser and R. E . Richards, ibid., 2, 357 (1959); R.Freeman, R. P. H. Gasser, and R. E. Richards, ibid., 2, 301 (1959). (3) J. H. E. George, J. A. Rolfe, and L. A. Woodward, Trans. Faraday Soc., 49, 375 (1953). (4) P. L. Goggins and L. A. Woodward, ibid., 66, 1591 (1960). ( 5 ) R. P. Bell and J. E. Prue, J . Chem. SOC.,362 (1949). ( 6 ) R. P. Bell and J. H. B. George, Trans. Faraday Soc., 49, 619 (1953). (7) R. P. Bell and M. H. Panekhurst. J . ‘Chem.SOC.,2836 (1956). (8) W. Ostwald, “Lehrbuch der Allseme.inen Chernie,” .Engelman, Leipzig, 1891. (9) M. Hlasko and A. Salitowna, Rocznilci Chem., 14, 1038 (1934); 16, 273 (1935).
Hlasko and Salitowna attempted to eliminate this interference by precipitation of BaCOs, their results evidently contain a systematic error, as indicated by the anomalous slope and curvature of the phoreogram and by comparison with the sum of ionic conductances a t infinite dilution. There is a corresponding systematic error in the Hlasko and Salitowna data on sodium hydroxide, potassium hydroxide, and lithium hydroxide, where comparison can be made with the more accurate results of Darken and Meierlo and Siverta, Reitmeier, and Tartar.” The purpose of the present work mas to obtain more accurate values for the equivalent conductance of dilute thallous hydroxide solutions a t 25’ by attempting to exclude carbon dioxide, and to determine the association constant from the results. Experimental Two series of experiments were undertaken. For eac!i series, fresh thallous hydroxide stock solutions, approximately 0.075 M , were prepared by bubbling carbon dioxidefree oxygen through de-ionized water in a closed system containing subdivided thallium metal which previously had been washed repeatedly with boiling de-ionized water. Spectrographic analyses of the subdivided metal showed the absence of any substantial concentrations of impurities. After dissolution of a major part of the metal, the resulting solution was decanted by gas pressure into a purged polyethylene storage bottle, where it wm kept in the dark under carbon dioxide-free nitrogen pressure. Provision was made for withdrawal of bolution from the storage bottle by hypodermic syringe. The stock solutions were analyzed a t several intervals during their use by weight potentiometric titrationb of triplicate samples against potassium acid phthalate primary standard. The TlOH solution was added from a hypodermic syringe which could be weighed rapidly and repeatedly by a projection-type analytical balance, while a stream of nitrogen was used to keep carbon dioxide (10) L. S. Darken and H. F. Meier, J . Am. Cham. Soc., 64, 621 (1942). (11) V. Sivertz, R. E. Reitrneier, and H. V . Tartar, ibid., 6 2 , 1379 (1940).
